This document provides information on potentiometry and potentiometric titration. It discusses the basic principles of potentiometry including electrode potentials and how a potential difference is established between an electrode and solution. It describes the instrumentation used including reference electrodes like calomel and silver-silver chloride electrodes and indicator electrodes like metal, glass membrane, and quinhydrone electrodes. It also discusses different types of potentiometric titrations and provides examples of applications for potentiometry in various industries.

1. Dr. B. R Ambedkar University
Department of Pharmacy,
Agra-282002
Guidance of Sir. Mr. Vijay Yadav

2. PRINCIPLE
 When a metal strip is placed in a solution
of its own ions there are two possibilities
or tendencies:
 Metal atoms may dissolve in the solution as
positive ions leaving electrons on the
electrode.
 Metal ions may take up electrons from
electrode and get deposited as neutral atoms
 In this way A POTENTIAL DIFFERENCE is
setup b/w electrode and solution .

3. POTENTIOMETER
 It is instrument used to determine the
potential differences between a reference
electrode and an indicator electrode. These
two electrodes form electro chemical cell that
are dipped in solution to be analyzed.
 Measured potential can be used to determine
the quantity of analyte in terms of
concentration.

5. ELECTRODE POTENTIAL
Tendency of electrode to lose or gain electrons
 Standard Potential
Potential of pure metal when it is dipped
in 1 Molar solution of its own ions at 25°C
(298K) is known as standard Electrode
potential.

6.  Oxidation Potential
The potential of substance to get oxidized is
called oxidation potential.
 Reduction Potential
The potential of substance to get reduced is
called reduction potential.

7. INSTRUMENTATION
 Potentiometric cell has following parts :
 Electrodes
 Salt bridge
 Analytical solution
 Galvanometer

9. WORKING
At its most basic, a potentiometer consists of two
electrodes, whose reduction potentials differ,
inserted in a test solution. The voltmeter is attached
to the electrodes to measure the potential difference
between them.
One of the electrodes is a reference electrode,
whose electrode potential is known.
The other electrode is the test electrode.
The test electrode is usually either a metal
immersed in a solution of its own ions, whose
concentration you wish to discover, or a carbon rod
electrode sitting a solution which contains the ions
of interest in two different oxidation states.

10. 1. Oxidation take place at anode.
For example on zinc electrode
Zn → Zn2+ + 2e-
1. Reduction take place at cathode.
For example on copper electrode
Cu2+ + 2e−→ Cu
When redox reaction take place than potential
is develop which is measured by
galvanometer.

11.  Electrode Potential is the potential of
electrochemical cell and can be represented
as :
E cell = E indicator - E reference
+E j
Ej= potential develops across the liquid
junction at each end of salt bridge it is
negligible .

13. ELECTRODES
Reference Electrode
1. Standard Hydrogen electrode
2. Calomel electrode
3. Silver-Silver Chloride electrode
Indicator Electrode
1. Metal Electrode
2. Hydrogen electrode
3. Quinhydrone electrode
4. Glass Membrane Electrode

14. REFERENCE
ELCTRODES

15. SHE
The standard H2 electrode potential is defined as the potential that is
developed between the H2 gas adsorbed on the pt metal and H+ of the
solution when the H2 gas at a pressure of 760 mm of Hg is in
equilibrium with H+ of unit concentration
Working:
 Pt foil : coated with black pt
 1 molar HCl solution
 Pure H2 gas bubbled continuously at 1 atm.
 Pt act as conductor, inert and facilitate in
attaining equilibrium
 Electrode Potential is 0.00 volts
 H2 (gas) ↔ 2H+ (ions) + 2e-
(electrons)

17.  Limitations :
1. Can’t be used in solution containing strong
oxidizing agents.
2. Difficult and expensive to maintain.
3. Excess of H2 bubbling out carries little HCl
with it and hence the H+ concentration
decreases. In such a system, it is difficult to
maintain the concentration of HCl at 1M.

18. CALOMEL ELECTRODE
 Most commonly used electrode in potentiometry
 Tube 5-15cm long, 0.5-1 cm in diameter.
 Slurry of mercury & mercurous chloride with
saturated soln of KCl
 Connected by a small opening with saturated
solution of KCl.
 Pt metal is placed inside the slurry
 Ceramic fiber act as salt bridge
 2Hg +2Cl¯ → Hg2Cl2 + 2e¯
 E = 0.2444 at 25°C

19. CALOMEL
ELECTRODE

20.  Advantages:
 Concentration of chloride ions don’t change
even some of the solvent get evaporated.
 Generates small junction potential so more
accurate
 Limitations :
 Mostly saturated solution of KCl is used and
that is temperature dependent.

21. SILVER-SILVER CHLORIDE ELECTRODE
 It consist of silver wire coated with AgCl
 Coating may be electroplating or Physical.
 This coated wire is placed in 1M solution of
AgCl.
 Ag + Cl¯ ↔ AgCl + 1e¯
 E= 0.199V
Advantages:
 Easy handling, and cost effective
 Limitation: sometimes show reactivity.

22. SILVER-SILVER CHLORIDE ELECTRODE

23. INDICATOR
ELECTRODES

24. METAL INDICATOR
ELECTRODESMetal electrode develops electric potential as a
result of redox reaction at its surface.
1. First Order Electrode
A first order electrode involves the metal in contact
with its own ions, such as Ag, Ag+ or Zn, Zn.
Ag⁺ + e¯ ↔ Ag (s) E= 0.800
Limitation:
 metallic indicator electrodes are not very selective
and
respond not only to their own cations but also to other
more easily reduced cations.
 – Many metal electrodes can be used only in neutral
or basic solutions because they dissolve in the
presence of acids

25. 2.SECOND ORDER ELECTRODE
A second-order electrode is one that responds to the
presence of precipitating or complexing ions. For
example, silver wire could serve as the indicator
electrode for chloride.
Ag +Cl¯ ↔ AgCl + e¯ E= 0.199

26. 3. INERT ELECTRODE
The most commonly used indicator electrodes
are known as inert. These electrodes are not
involved in the half-cell reactions of the
electrochemical species. Typical inert
electrodes are platinum, gold, and carbon.
Inert electrodes are responsive to any
reversible redox system; these are widely
used in potentiometric work.

27. GLASS MEMBRANE ELECTRODE
 Its most commonly used indicator electrode
 It involves Ion Exchange reaction
 Membrane is made up of chemically bonded
Na2O, SiO2 and Al2O3
 The Glass bulb is filled with solution of HCl &
KCl, silver acetate coated with AgCl is
inserted as Electrode.

30. WORKING
 First glass absorb water
 Then H ions can move in the direction of lesser
concentration and replace Na ions and others in
the glass membrane.
 SiO2-Na⁺ + H⁺ → SiO- H⁺ + Na⁺
 As a result of diffusion and exchange process a
potential develops on each side of glass membrane

31. APPLICATIONS
 Its potential is not effected by the presence of
strong reducing and oxidizing agents.
 It operates over a wide pH range,
 It respond fast and function well in
physiological systems
 Used in pH meter.

33. HYDROGEN ELECTRODE
 Used in acid base titration but very limited
application because many organic
substances directly react with hydrogen
gas.

34. QUINHYDRONE ELECTRODE
 Contain a solution of Quinones and
Hydroquinones prepared from Quinhydrone
 2QH → Q+H2Q
 Simple and easily attains equilibrium
 employed in the presence of mild oxi and
red agents at pH8
 Sensitive to high conc of salts and oxi & red
agents.

35.  The electrode consists of an inert metal
electrode (usually a platinum wire) in contact
with quinhydrone crystals and a water-based
solution. Quinhydrone is slightly soluble in
water, dissolving to form a mixture of two
substances, quinone and hydroquinone, with
the two substances present at equal
concentration. Each one of the two substances
can easily be oxidised or reduced to the other.
 The potential at the inert electrode depends on
the ratio of the activity of two substances
(quinone-hydroquinone), and also the
hydrogen ion concentration. The electrode half-
reaction is:
 Hydroquinone ↔ Quinone + 2H+ +2e-

36.  For practical pH measurement, a second pH independent
reference electrode (such as a silver chloride electrode) is
also used. This reference electrode does not respond to
the pH. The difference between the potential of the two
electrodes depends (primarily) on the activity of H+ in the
solution. It is this potential difference which is measured
and converted to a pH value.
 The quinhydrone electrode is not reliable above pH 8. It is
also unreliable in the presence of strong oxidising or
reducing agents, which would disturb the equilibrium
between hydroquinone and quinone. It is also subject to
errors in solutions containing proteins or high
concentrations of salts.
 Other electrodes commonly used for measuring pH are
the glass electrode, the hydrogen electrode, the antimony
– antimony oxide electrode, and the ion-sensitive field
effect transistor ISFET electrode.

37. SALT BRIDGE
 U shaped tube filled with an inert
electrolyte
1. Glass tube bridge (gel+ KI or Na2SO4)
2. Filter tube bridge (filter paper+KCl or NaCl)
Function:
1. Allow electrical contact b/w 2 solutions.
2. Prevent mixing of two solutions.
3. Maintain electrical neutrality.

39. TITRATION :
 Quantitative measuring procedure in which a
liquid solution is added to a mixture until
some distinctive feature, signal or end point
is observed.

40. POTENTIOMETRIC TITRATION
 It’s a volumetric method in which potential
between two electrodes (reference &
indicator) is measured as a function of added
reagent volume.

43. TYPES
1. Precipitation titration
2. Complex formation Titration
3. Neutralization titration
4. Oxidation – reduction titration

44. PRECIPITATION TITRATION
 Volumetric methods based upon the
formation of slightly soluble precipitate are
called " precipitation titration " .
 Because of the precipitating titration based
upon utilizing silver nitrate
 (AgNO3) as a precipitating agent, then
it called " argentimetric
 processes " .
 Precipitation titration is a very important ,
because it is a perfect method
 for determine halogens and some metal
ions

45. COMPLEX FORMATION TITRATION
 Complex ions ( coordination compounds) are
produce from reaction of
 many metal ions (electrons accepter) with
electron pair donors .
 The donor species (or called ligands) must
have at least one pair of
 unshared electrons for bond formation

46. NEUTRALIZATION TITRATION
 Neutralization titrations are performed
with standard solutions of strong acids
or
 bases. While a single solution (of either
acid or base) is sufficient for the titration
of
 a given type of analyte, it is convenient
to have standard solutions of both acid
and
 base available in case back-titration is
needed to locate the end point more

47. OXIDATION-REDUCTION TITRATION
 A redox titration is a type of titration based on a
redox reaction between the analyte and titrant.
Redox titration may involve the use of a redox
indicator and/or a potentiometer. Common
examples of a redox titration is treating a
solution of iodine with a reducing agent and
using starch as an indicator. Iodine forms an
intensely blue complex with starch. Iodine (I2)
can be reduced to iodide(I−) by
e.g. thiosulphate (S2O3
2−), and when all iodine is
spent the blue colour disappears. This is called
an iodometric titration

48. APPLICATIONS OF POTENTIOMETRY
 Analysis of pollutants in water
 Drug Analysis in Pharmaceutical industry
 Food industry for analysis of quality
 Biochemical and biological Assay or analysis
To check the quality of cosmetics
 Also used as analytical tool in Textile, paper,
paints, explosive energy and more .